Write the Lewis structure of the following compounds and show the formal charge on each atom: $HNO_3, NO_2, H_2SO_4$.

(N/A) The Lewis structures and formal charges on each atom are calculated using the formula: $\text{Formal charge} = [\text{Total valence electrons}] - [\text{Total non-bonding electrons}] - \frac{1}{2} [\text{Total shared electrons}]$.
$(i)$ $HNO_3$:
Formal charge on $H = 1 - 0 - \frac{1}{2}(2) = 0$
Formal charge on $N = 5 - 0 - \frac{1}{2}(8) = +1$
Formal charge on $O(1) = 6 - 4 - \frac{1}{2}(4) = 0$
Formal charge on $O(2) = 6 - 4 - \frac{1}{2}(4) = 0$
Formal charge on $O(3) = 6 - 6 - \frac{1}{2}(2) = -1$
$(ii)$ $NO_2$:
Formal charge on $O(1) = 6 - 4 - \frac{1}{2}(4) = 0$
Formal charge on $N = 5 - 1 - \frac{1}{2}(6) = +1$
Formal charge on $O(2) = 6 - 6 - \frac{1}{2}(2) = -1$
$(iii)$ $H_2SO_4$:
In $H_2SO_4$,the central $S$ atom is bonded to two $OH$ groups and two $O$ atoms via double bonds. All atoms follow the octet rule (except $H$). The formal charge on $S = 6 - 0 - \frac{1}{2}(12) = 0$. The formal charge on $OH$ oxygen atoms $= 6 - 4 - \frac{1}{2}(4) = 0$. The formal charge on double-bonded oxygen atoms $= 6 - 4 - \frac{1}{2}(4) = 0$.